1. Stoichiometry: Calculations with Chemical Formulas and Equations
Chapter 3

2. An equation Describes a reaction
Must be balanced to follow Law of Conservation of Energy
Can only be balanced by changing the coefficients.
Has special symbols to indicate state, and if catalyst or energy is required.

3. Symbols used in equations
the arrow separates the reactants from the products.
(s) after the formula –solid
(g) after the formula –gas
(l) after the formula –liquid
(aq) after the formula-aqueous (Dissolved in water, and aqueous solution)

4. Symbols used in equations
indicates a reversible reaction
shows that heat is supplied to the reaction
is used to indicate a catalyst used supplied, in this case, platinum.

5. What is a catalyst?
A substance that speeds up a reaction without being changed by the reaction.
ex.. enzymes-biological catalysis
protein catalysts

6. Convert these to equations
Solid iron (II) sulfide reacts with gaseous hydrogen chloride to form iron (II) chloride and hydrogen sulfide gas.
Nitric acid dissolved in water reacts with solid sodium carbonate to form liquid water and carbon dioxide gas and sodium nitrate dissolved in water.

7. Balancing Chemical Equations

8. Balanced Equation Atoms can’t be created or destroyed
A balanced equation has the same number of each element on both sides of the equation.

9. Rules for balancing Equations
Write the correct formulas for all the reactants and products
Count the number of atoms of each type appearing on both sides
Balance the elements one at a time by adding coefficients (the numbers in front)
Check to make sure it is balanced.

10. Never Change a subscript to balance an equation
If you change the formula you are describing a different reaction
H2O is a different compound than H2O2
Never put a coefficient in the middle of a formula
2 NaCl is okay, Na2Cl is not.

11. Balance the following equation
AgNO3 + Cu → Cu(NO3)2 + Ag
Mg + N2 → Mg3N2
P + O2 → P4O10
Na + H2O → H2 + NaOH
CH4 + O2 → CO2 + H2O

12. Predicting The products
Types of Reactions
Predicting The products

13. Types of Reactions There are millions of reactions.
Fall into several categories.
Combination
Decomposition
Single Replacement
Double Displacement
Combustion

14. Combination Reactions
2 elements, or compounds combine to make one compound.
Ca +O2 ® CaO
SO3 + H2O ® H2SO4
We can predict the products if they are two elements.
Mg + N2 ®

15. Write and balance Ca + Cl2 → Fe + O2 →iron (II) oxide Al + O2 →
Remember that the first step is to write the formula
Then balance

16. Decomposition Reactions
decompose = fall apart
one reactant falls apart into two or more elements or compounds.
NaCl Na + Cl2
CaCO CaO + CO2

17. Double Replacement Two things replace each other.
Reactants must be two ionic compounds or acids.
Usually in aqueous solution
NaOH + FeCl3 ®
The positive ions change place.
NaOH + FeCl3 ® Fe+3 OH- + Na+1Cl-1
NaOH + FeCl3 ® Fe(OH)3 + NaCl

18. Double Replacement Will only happen if one of the products
doesn’t dissolve in water and forms a solid
or is a gas that bubbles out.
or is a covalent compound usually water.

19. Complete and balance assume all of the reactions take place.
CaCl2 + NaOH ®
CuCl2 + K2S ®
KOH + Fe(NO3)3 ®
(NH4)2SO4 + BaF2 ®

20. How to recognize which type
Look at the reactants
E + E Combination
C Decomposition
C + C Double replacement

21. Combustion
A compound composed of only C H and maybe O is reacted with oxygen
If the combustion is complete, the products will be CO2 and H2O.
If the combustion is incomplete, the products will be CO and H2O.

22. Reactions Come in 5 types. (We learned 4)
Can tell what type they are by the reactants.
Double Replacement happens if the product is a solid, water, or a gas.

23. The Process Determine the type by looking at the reactants.
Put the pieces next to each other
Use charges to write the formulas
Use coefficients to balance the equation.

24. Chemical Quantities

25. Moles
Defined as the number of carbon atoms in exactly 12 grams of carbon-12.
1 mole is 6.02 x particles/molecules.
Treat it like a very large dozen
6.02 x is called Avogadro's number.

26. Representative particles
The smallest pieces of a substance.
For a molecular compound it is a molecule.
For an ionic compound it is a formula unit.
For an element it is an atom.

27. Types of questions How many oxygen atoms in the following?
CaCO3
Al2(SO4)3
How many ions in the following?
CaCl2
NaOH

28. Types of questions
How many molecules of CO2 are the in 4.56 moles of CO2 ?
How many moles of water is 5.87 x molecules?
How many atoms of carbon are there in 1.23 moles of C6H12O6 ?
How many moles is 7.78 x 1024 formula units of MgCl2?

29. Measuring Moles The amu was one twelfth the mass of a carbon 12 atom.
Since the mole is the number of atoms in 12 grams of carbon-12,
the decimal number on the periodic table is also the mass of 1 mole of those atoms in grams.

30. Gram Atomic Mass The mass of 1 mole of an element in grams.
12.01 grams of carbon has the same number of pieces as grams of hydrogen and grams of iron.
We can write this as g C = 1 mole

31. Examples How much would 2.34 moles of carbon weigh?
How many moles of magnesium in g of Mg?
How many atoms of lithium in 1.00 g of Li?
How much would 3.45 x 1022 atoms of U weigh?

32. What about compounds?
in 1 mole of H2O molecules there are two moles of H atoms and 1 mole of O atoms
To find the mass of one mole of a compound
determine the moles of the elements they have
Find out how much they would weigh
add them up

33. What about compounds? What is the mass of one mole of CH4?
1 mole of C = g
4 mole of H x 1.01 g = 4.04g
1 mole CH4 = = 16.05g
The Gram Molecular mass of CH4 is 16.05g
The mass of one mole of a molecular compound.

34. Atomic Mass The mass of one mole of an ionic compound.
Calculated the same way.
What is the atomic mass of Fe2O3?
2 moles of Fe x g = g
3 moles of O x g = g
The atomic mass = g g = g

35. Molar Mass The generic term for the mass of one mole.
The same as gram molecular mass, gram formula mass, and gram atomic mass.

36. Examples Calculate the molar mass of the following. Na2S N2O4 C
Ca(NO3)2
C6H12O6
(NH4)3PO4

37. Using Molar Mass Finding moles of compounds
Counting pieces by weighing

38. Molar Mass The number of grams of 1 mole of atoms, ions, or molecules.
We can make conversion factors from these.
To change grams of a compound to moles of a compound.

39. For example
How many moles is 5.69 g of NaOH?

40. For example
How many moles is 5.69 g of NaOH?

41. For example How many moles is 5.69 g of NaOH?
need to change grams to moles
for NaOH
1mole Na = 22.99g 1 mol O = g 1 mole of H = 1.01 g
1 mole NaOH = g

42. Examples How many moles is 4.56 g of CO2 ?
How many grams is 9.87 moles of H2O?
How many molecules in 6.8 g of CH4?
49 molecules of C6H12O6 weighs how much?

43. Mole to Mole conversions
How many moles of O2 are produced when 3.34 moles of Al2O3 decompose?
2 Al2O3 ® 4Al + 3O2
3.34 moles Al2O3
3 mole O2 =
5.01 moles O2
2 moles Al2O3

44. Your Turn
2C2H2 + 5 O2 ® 4CO2 + 2 H2O
If 3.84 moles of C2H2 are burned, how many moles of O2 are needed?
How many moles of C2H2 are needed to produce 8.95 mole of H2O?
If 2.47 moles of C2H2 are burned, how many moles of CO2 are formed?

45. How do you get good at this?

46. Mass in Chemical Reactions
How much do you make?
How much do you need?

47. For example...
If 10.1 g of Fe are added to a solution of Copper (II) Sulfate, how much solid copper would form?
Fe + CuSO4 ® Fe2(SO4)3 + Cu
2Fe + 3CuSO4 ® Fe2(SO4)3 + Cu
1 mol Fe
10.1 g Fe =
0.181 mol Fe
55.85 g Fe

48. 2Fe + 3CuSO4 ® Fe2(SO4)3 + 3Cu
3 mol Cu
0.272 mol Cu
0.181 mol Fe =
2 mol Fe
63.55 g Cu
0.272 mol Cu =
17.3 g Cu
1 mol Cu

49. Could have done it 1 mol Fe 63.55 g Cu 10.1 g Fe 3 mol Cu 55.85 g Fe=
17.3 g Cu

50. More Examples
To make silicon for computer chips they use this reaction
SiCl4 + 2Mg ® 2MgCl2 + Si
How many grams of Mg are needed to make 9.3 g of Si?
How many grams of SiCl4 are needed to make 9.3 g of Si?
How many grams of MgCl2 are produced along with 9.3 g of silicon?

51. For Example The U. S. Space Shuttle boosters use this reaction
3 Al(s) + 3 NH4ClO4 ® Al2O3 + AlCl3 + 3 NO + 6H2O
How much Al must be used to react with 652 g of NH4ClO4 ?
How much water is produced?
How much AlCl3?

52. Gas and Moles

53. Gases Many of the chemicals we deal with are gases.
They are difficult to weigh.
Need to know how many moles of gas we have.
Two things effect the volume of a gas
Temperature and pressure

54. Standard Temperature and Pressure (STP)
0ºC and 1 atm pressure
At STP 1 mole of gas occupies 22.4 L
Called the molar volume
Avogadro's Hypothesis - at the same temperature and pressure equal volumes of gas have the same number of particles.

55. For Example
If 6.45 grams of water are decomposed, how many liters of oxygen will be produced at STP?
2H2O ® 2H2 + O2
1 mol H2O
1 mol O2
22.4 L O2
6.45 g H2O
18.02 g H2O
2 mol H2O
1 mol O2

56. Example
How many liters of CO2 at STP will be produced from the complete combustion of 23.2 g C4H10 ?
What volume of oxygen will be required?

57. Example
How many liters of CH4 at STP are required to completely react with 17.5 L of O2 ?
CH4 + 2O2 ® CO2 + 2H2O
22.4 L O2
1 mol O2
1 mol CH4
22.4 L CH4
1 mol O2
1 mol CH4
22.4 L CH4
17.5 L O2
22.4 L O2
2 mol O2
1 mol CH4
= 8.75 L CH4

58. Density of a gas D = m /V for a gas the units will be g / L
We can determine the density of any gas at STP if we know its formula.
To find the density we need the mass and the volume.
If you assume you have 1 mole than the mass is the molar mass (PT)
At STP the volume is 22.4 L.

59. Examples Find the density of CO2 at STP.
Find the density of CH4 at STP.

60. The other way
Given the density, we can find the molar mass of the gas.
Again, pretend you have a mole at STP, so V = 22.4 L.
m = D x V
m is the mass of 1 mole, since you have 22.4 L of the stuff.
What is the molar mass of a gas with a density of g/L?
2.86 g/L?

61. Limiting Reagent

62. Limiting Reagent
If you are given one dozen loaves of bread, a gallon of mustard and three pieces of salami, how many salami sandwiches can you make (rhetorical question).
The limiting reagent is the reactant you run out of first.
The excess reagent is the one you have left over.
The limiting reagent determines how much product you can make

63. How do you find out? Do two stoichiometry problems.
The one that makes the least product is the limiting reagent.
For example
Copper reacts with sulfur to form copper ( I ) sulfide. If 10.6 g of copper reacts with 3.83 g S how much product will be formed?

64. If 10. 6 g of copper reacts with 3. 83 g S
If 10.6 g of copper reacts with 3.83 g S. How many grams of product will be formed?
2Cu + S ® Cu2S
Cu is Limiting Reagent
1 mol Cu
1 mol Cu2S
g Cu2S
10.6 g Cu
63.55g Cu
2 mol Cu
1 mol Cu2S
= 13.3 g Cu2S
= 13.3 g Cu2S
1 mol S
1 mol Cu2S
g Cu2S
3.83 g S
32.06g S
1 mol S
1 mol Cu2S
= 19.0 g Cu2S

65. Your turn
If 10.1 g of magnesium and 2.87 L of HCl gas are reacted, how many liters of gas will be produced?
How many grams of solid?
How much excess reagent remains?

66. Your Turn II
If 10.3 g of aluminum are reacted with 51.7 g of CuSO4 how much copper will be produced?
How much excess reagent will remain?

67. Yield The amount of product made in a chemical reaction.
There are three types
Actual yield- what you get in the lab when the chemicals are mixed
Theoretical yield- what the balanced equation tells you you should make.
Percent yield = Actual x 100 % Theoretical

68. Example
6.78 g of copper is produced when 3.92 g of Al are reacted with excess copper (II) sulfate.
2Al + 3 CuSO4 ® Al2(SO4)3 + 3Cu
What is the actual yield?
What is the theoretical yield?
What is the percent yield?

69. Details Percent yield tells us how “efficient” a reaction is.
Percent yield can not be bigger than 100 %.

70. Empirical Formula
From percentage to formula

71. The Empirical Formula
The lowest whole number ratio of elements in a compound.
The molecular formula the actual ratio of elements in a compound.
The two can be the same.
CH2 empirical formula
C2H4 molecular formula
C3H6 molecular formula
H2O both

72. Calculating Empirical
Just find the lowest whole number ratio
C6H12O6
CH4N
It is not just the ratio of atoms, it is also the ratio of moles of atoms.
In 1 mole of CO2 there is 1 mole of carbon and 2 moles of oxygen.
In one molecule of CO2 there is 1 atom of C and 2 atoms of O.

73. Calculating Empirical
We can get ratio from percent composition.
Assume you have a 100 g.
The percentages become grams.
Can turn grams to moles.
Find lowest whole number ratio by dividing by the smallest.

74. Example
Calculate the empirical formula of a compound composed of % C, % H, and %N.
Assume 100 g so
38.67 g C x 1mol C = mole C gC
16.22 g H x 1mol H = mole H gH
45.11 g N x 1mol N = mole N gN

75. Example The ratio is 3.220 mol C = 1 mol C 3.219 mol N 1 mol N
The ratio is mol H = 5 mol H mol N mol N
C1H5N1
A compound is % P and % O. What is the empirical formula?
Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. What is its empirical formula?

76. Empirical to molecular
Since the empirical formula is the lowest ratio the actual molecule would weigh more.
By a whole number multiple.
Divide the actual molar mass by the the mass of one mole of the empirical formula.
Caffeine has a molar mass of 194 g. what is its molecular mass?

77. Example
A compound is known to be composed of % Cl, 24.27% C and 4.07% H. Its molar mass is known (from gas density) is known to be g. What is its molecular formula?